Intermolecular Forces

Intermolecular Forces

• Intermolecular forces arise from the charged nature of subatomic particles (electrons and protons).

• They decrease with increasing average particle separation.

• Intermolecular forces are responsible for the cohesiveness of materials.

• They determine the physical properties of pure substances such as melting point, boiling point, and vapor pressure.

Intermolecular Forces vs. Phase

• Substances that are gases at room temperature have weak intermolecular forces.

• Substances that are condensed (liquids or solids) at room temperature have much stronger intermolecular forces.

• If intermolecular forces did not exist, all substances would be gases, even at extremely low temperatures.

Intermolecular Forces Determine Physical State

• The fundamental difference between states of matter is the distance between particles.

Condensed Phases

• The solid and liquid states are referred to as condensed phases because the particles are closer together.

Intermolecular Forces and Kinetic Energy

• The physical state of a substance at a particular temperature and pressure depends on two opposing entities:

  • The kinetic energy of the particles.

  • The strength of the intermolecular forces (potential energy) between the particles.

• These forces are generally the attractive ones between particles in a solid, liquid, or gas.

• Converting from one physical state to another requires the molecules to gain enough energy to overcome the intermolecular forces.

Types of Intermolecular Forces

• Ion-dipole (strongest)

  • Applies to solutions, not pure substances.

  • The solvent must be polar.

• Hydrogen bonding

  • Substance must be polar.

• Dipole-dipole

  • Substance must be polar.

• Dispersion (weakest)

  • Substance must contain both protons and electrons.

The Effect of Intermolecular Forces on Boiling Point and Melting Point

• As the strength of intermolecular forces increases:

  • Boiling point (BP) increases.

  • Heat of vaporization (ar{ ext{vaph}}) increases.

  • Vapor pressure decreases.

  • Vapor pressure (VP) is the pressure exerted by a gas in equilibrium with its liquid phase.

  • VP and BP are inversely related.

  • Melting point (MP) increases.

  • Heat of fusion (ar{ ext{fus}}) increases.

Dipole-Dipole Force

• Dipole-dipole interactions (forces) are found between polar molecules.

• These forces result from the attraction (or repulsion) of the positive and negative ends of the dipole moments of polar molecules.

• The most stable arrangement of polar molecules has the positive end of one molecule oriented toward the negative end of another molecule.

Example of Dipole-Dipole Force

• The largest intermolecular force between molecules of CH₃Cl comes from the dipole-dipole interaction.

Heat of Vaporization and Boiling Point

• In order for vaporization to occur, intermolecular forces must be overcome.

• As polarity increases, the strength of dipole-dipole interactions increases, thus requiring more energy for vaporization.

• Larger heat of vaporization (ar{ ext{vaph}}) results in higher boiling point (BP).

London Dispersion Forces

• London dispersion forces are present between all molecules.

• Temporary dipole moments last for fractions of a second, induced in one molecule by other nearby molecules, leading to displacement of electrons from a symmetrical arrangement.

Strength of the London Dispersion Force Increases with Surface Area

• London dispersion forces require close surface contact between two (or more) molecules.

• The strength of dispersion forces is roughly proportional to the surface area in contact between the molecules.

• As surface area increases, London dispersion forces increase, leading to higher boiling points.

Dispersion Force and Boiling Point

• For molecules with similar molecular formulas:

  • Long, skinny (unbranched) molecules have greater surface area, resulting in higher boiling points.

  • Shorter, branched molecules are more spherical and have less surface area, resulting in lower boiling points.

Strength of the London Dispersion Force Tends to Increase with Increasing Molecular Weight

• Larger atoms have larger electron clouds, which are easier to polarize.

• In general, as molecular weight (MW) increases, boiling point (BP) increases.

The Boiling Point of Water

• Although boiling point tends to increase with increasing molecular weight, there are exceptions.

• The boiling point of water is higher than expected due to strong hydrogen bonding.

Hydrogen Bonding

• Compounds containing H-F, O-H, and/or N-H bonds exhibit hydrogen bonding:

  • Strong dipole-dipole interaction between a hydrogen atom covalently bonded to O, N, or F and a lone pair of electrons on a different O, N, or F atom.

Hydrogen Bonding is a Strong Intermolecular Force

• Due to the large differences in electronegativity, O-H, N-H, and F-H bonds are highly polar.

• The hydrogen atom has a strong partial positive charge, leading to a strong attraction to the nonbonding electrons on other N, O, or F atoms.

Identifying Molecules That Can Hydrogen Bond

• The ability to form hydrogen bonds depends on the presence of the appropriate bonds in the structure of a molecule.

H-Bonding and Boiling Point

• Impact of hydrogen bonding on boiling point (BP):

  • Hydrogen bonding leads to higher boiling points.

  • Water (H₂O) forms hydrogen bonds; H₂S cannot form H-bonds.

  • As the number of hydrogens capable of forming hydrogen bonds increases, the boiling point increases.

H-Bonding is Stronger with OH than with NH

• Due to greater differences in electronegativity, OH forms stronger hydrogen bonds than NH.

• Compounds with OH have higher boiling points than similar compounds with NH or NH₂. (e.g., BP of CH₃CH₂CH₂CH₂NH₂ = 77-78 °C vs. BP of CH₃CH₂CH₂CH₂OH = 117-118 °C)

Predict and Explain Relative BPs - I

• Comparison of pairs of compounds for boiling point:

  • Identify which compound has a higher BP and explain the reasoning.

  • Evaluate relative vapor pressures.

Predict and Explain Relative BPs - II

• Similar analysis of boiling points for different compound pairs, considering influence of molecular structure on boiling points.

Solubility and Polarity

• Polar liquids tend to dissolve readily in polar solvents due to dipole-dipole interactions.

  • Example: Acetone and water are both polar and miscible.

• Nonpolar liquids tend to be insoluble in polar liquids.

  • Example: Gasoline and water are immiscible.

Solubility – Like Dissolves Like

• Nonpolar liquids tend to dissolve readily in nonpolar solvents due to dispersion forces.

  • Example: Gasoline and hexane are both nonpolar and miscible.

• The general rule is “like dissolves like.”

Intermolecular Forces and Solubility

• Polar compounds dissolve in polar solvents due to:

  • Dipole-dipole interactions.

  • Hydrogen bonding (sometimes).

  • Example of hydrogen bonding: between methyl amine and water.

Why Nonpolar Compounds Do Not Dissolve Appreciably in Water

• Nonpolar compounds do not dissolve appreciably in water because they cannot break the hydrogen-bonding network that exists in water.

Solubility and the Ion-Dipole Force

• Ionic compounds tend to dissolve in polar solvents like water due to ion-dipole forces.

• Ion-dipole attractions cannot occur in a pure substance; they are the attractive force between an ion and the partially charged end of a polar molecule.

Solvation of Ions in Water

• An aqueous solution of an ionic compound such as NaCl contains solvated cations (C Na⁺) and anions (C Cl⁻).

• This solvation process highlights examples of ion-dipole forces.

Predicting Solubility in Water

• Evaluate solubility of various solutes in water:

  • CH₃CH₂CH₃ vs. CH₃CH₂OH

  • HOCH₂CH₂CH₂CH₂OH vs. CH₃CH₂CH₂CH₂CH₂OH

  • Vitamin A vs. Vitamin C

Solubility of Large Molecules

• When both solute and solvent contain polar and nonpolar components:

  • Estimate solubility by assessing the polar and nonpolar portions.

• As the length of the carbon chain in alcohols (R-OH) increases, solubility in water decreases, while solubility in hexane (C₆H₁₄) increases.